FORCES IN SOLIDS AND
LIQUIDS
(INTERMOLECULAR FORCES)
IONIC BOND
METALLIC BOND
COVALENT BOND
VAN DER WAALS FORCES
DIPOLE-DIPOLE ATTRACTIONS
HYDROGEN BONDING
INTERMOLECULAR FORCES
Introduction:
The physical properties of melting point,
boiling point, vapor pressure, evaporation,
viscosity, surface tension, and solubility are
related to the strength of attractive forces
between molecules.
These attractive forces are called
Intermolecular Forces.
1. IONIC FORCES
The forces holding ions together in ionic
solids are electrostatic forces.
Opposite charges attract each other.
These are VERY strong intermolecular
forces.
Ionic forces hold many ions in a crystal
lattice structure.
IONIC FORCES
A small representative bit of a sodium
chloride lattice
exploded version of a sodium chloride
lattice
Crystalline Lattice structure of
NaCl(s)
Ionic Bonding
LATTİCE ENERGY is the energy
gained in converting 1 mol of solid
NaCl into its gaseous ions.
Lattice Energy
Depends on;
1) Ionic charge
Ionic charge
lattice enrgy
Ionic bond strength
Melting point of the solid
Lattice Energy
Depends on;
2) Atomic radii
Atomic radii
lattice energy
melting point of the solid
Ionic bond strength
Ionic compounds
Are solid at room temperature
Have crystalline structure
Have high melting and boiling points
Do not conduct the electricity in their
solid state
Can only conduct electricity in their
liquid state and aqueous solutions
Are not ductile or malleable
METALLIC FORCE
Metals tend to have high melting points and
boiling points suggesting strong bonds
between the atoms.
Even a metal like sodium (melting point
97.8°C) melts at a considerably higher
temperature than the element (neon) which
precedes it in the Periodic Table.
METALLIC BOND
electron sea
Metallic bonding in sodium
The electrons can
move freely within
these molecular
orbitals, and so each
electron becomes
detached from its
parent atom. The
electrons are said to
be delocalised. The
metal is held together
by the strong forces of
attraction between the
positive nuclei and the
delocalised electrons.
Metallic bonding in sodium
This is sometimes described as
"an array of positive ions in a
sea of electrons".
If you are going to use this view,
beware! Is a metal made up of
atoms or ions? It is made of
atoms.
Each positive centre in the
diagram represents all the rest of
the atom apart from the outer
electron, but that electron hasn't
been lost - it may no longer have
an attachment to a particular
atom, but it's still there in the
structure. Sodium metal is
therefore written as Na - not Na+
or potassium metal is written as
K but not K+.
Metallic bond strength
Depends on
1) Valence electron number of
the metal
2) Atomic radii of the metal
Metallic bonding in magnesium
If you work through the same argument with magnesium,
you end up with stronger bonds and so a higher melting
point.
Magnesium has the outer electronic structure 3s2. Both of
these electrons become delocalised, so the "sea" has
twice the electron density as it does in sodium. The
remaining "ions" also have twice the charge (if you are
going to use this particular view of the metal bond) and so
there will be more attraction between "ions" and "sea".
More realistically, each magnesium atom has one more
proton in the nucleus than a sodium atom has, and so not
only will there be a greater number of delocalised
electrons, but there will also be a greater attraction for
them.
Magnesium atoms have a slightly smaller radius than
sodium atoms, and so the delocalised electrons are closer
to the nuclei. Each magnesium atom also has twelve near
neighbours rather than sodium's eight. Both of these
factors increase the strength of the bond still further.
In periodic table;
Metallic bond strength increases
Metallic bond strength
decreases.
The metallic bond in molten
metals
In a molten metal, the metallic bond is
still present, although the ordered
structure has been broken down.
The metallic bond isn't fully broken
until the metal boils. That means that
boiling point is actually a better guide
to the strength of the metallic bond
than melting point is.
On melting, the bond is loosened, not
broken.
The physical properties of metals
Melting points and boiling points
Metals tend to have high melting and boiling points
because of the strength of the metallic bond. The
strength of the bond varies from metal to metal and
depends on the number of electrons which each atom
delocalises into the sea of electrons, and on the
packing.
Group 1 metals like sodium and potassium have
relatively low melting and boiling points mainly because
each atom only has one electron to contribute to the
bond.
They have relatively large atoms (meaning that the
nuclei are some distance from the delocalised
electrons) which also weakens the bond.
The physical properties of metals
Electrical conductivity
Metals conduct electricity. The
delocalised electrons are free to
move
Liquid metals also conduct
electricity, showing that although
the metal atoms may be free to
move, the delocalisation remains
in force until the metal boils.
Electrical conductivity of metals
Animation
showing
electrons
moving
randomly and
then the
movement of
electrons
through a wire
Electrical conductivity of
metals
Their electrical conductivity
decreases as temperature
increases.
Thermal conductivity
Metals are good conductors of heat.
Heat energy is picked up by the
electrons as additional kinetic energy
(it makes them move faster).
The energy is transferred throughout
the rest of the metal by the moving
electrons.
Thermal conductivity of metals
Metals are good conductors of heat. There
are two reasons for this:
the close packing of the metal ions in the
lattice
the delocalised electrons can carry kinetic
energy through the lattice
Strength and workability
Malleability and ductility
Metals are described as malleable (can be beaten into sheets) and
ductile (can be pulled out into wires).
This is because of the ability of the atoms to roll over each other
into new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will start
to roll over each other.
If the stress is released again, they will fall back to their original
positions. Under these circumstances, the metal is said to be
elastic.
Because of the electron sea,
metals have lustrous
appearance.
Strength and workability
If a larger stress is put on, the atoms
roll over each other into a new position,
and the metal is permanently changed.
As metallic bond strength
increases, the metallic activity
decreases.
Metallic Bonding
A. Outermost electrons wander freely through metal. Metal consists of
cations held together by negatively-charged electron "glue.“
B. Free electrons can move rapidly in response to electric fields,
hence metals are a good conductor of electricity.
C. Free electrons can transmit kinetic energy rapidly, hence metals are
good conductors of heat.
D. The layers of atoms in metal are hard to pull apart because of the
electrons holding them together, hence metals are tough. But
individual atoms are not held to any other specific atoms, hence
atoms slip easily past one another. Thus metals are ductile. Metallic
Bonding is the basis of our industrial civilization.
2. DIPOLE FORCES
Polar covalent molecules are sometimes
described as "dipoles", meaning that the
molecule has two "poles". One end
(pole) of the molecule has a partial
positive charge while the other end has a
partial negative charge. The molecules
will orientate themselves so that the
opposite charges attract principle
operates effectively.
FORCES BETWEEN MOLECULES
There are in fact three basic types of
interaction between molecules which
are, in order of increasing strength:
van der Waals interactions or
dispersion forces,
dipole-dipole interactions,
hydrogen bonds,
and these secondary bonding
interactions, like the primary bonding
interactions, involve the electrons
(atomic glue!).
3. VAN DER WAALS (LONDON)
FORCES:
3. VAN DER WAALS (LONDON)
FORCES:
All molecules have these forces.
However some molecules have ONLY
Van Der Waals forces but no other
forces in their solid and liquid states.
These are;
1) Noble gas atoms
2) Nonpolar molecules (CH4, CO2, C2H2,
etc.)
3. VAN DER WAALS (LONDON)
FORCES:
Van Der Waals forces appear on
the surface of the molecules.
Therefore, as the surface area of
the molecule increases, the
strength of Van Der Waals forces
also increases.
3. VAN DER WAALS (LONDON)
FORCES:
As the total # of electrons increases in the
molecule, the size of the molecule get larger
and the strength of Van Der Walls forces
increases.
The straight chain molecules have stronger
Van Der Waals forces than the branched
molecules.
3. VAN DER WAALS (LONDON)
FORCES:
n-pentane (boiling point:
36°C) has stronger Van Der
neopentane
(boiling
Waals forces than
point: 9°C)
neopentane
3. VAN DER WAALS (LONDON)
FORCES:
H2 has 2 e
F2 has 18 e
F2 has stronger Van Der Waals
forces than H2, and therefore
higher boiling point than H2.
3. VAN DER WAALS (LONDON)
FORCES:
Are the weakest forces!!!!
Therefore, the molecules
having these forces only will
always have the lowest melting
or boiling points.
4. DIPOLE FORCES
In the example on the
right,hydrochloric acid
is a polar molecule
with the partial positive
charge on the
hydrogen and the
partial negative charge
on the chlorine.
A network of partial +
and - charges attract
molecules to each
other.
DIPOLE FORCES
Happen between polar
molecules.
DIPOLE FORCES
Appear on the surface of the
molecules. Therefore, their strength
depends on the size (total e number) of
the molecule.
MOLECULES HAVING DIPOLE-DIPOLE
ATTRACTIONS WILL ALSO HAVE VAN
DER WAALS FORCES!!!
5. HYDROGEN BONDING
The hydrogen bond is really a special
case of dipole forces. A hydrogen bond is
the attractive force between the hydrogen
attached to an electronegative atom of
one molecule and an electronegative
atom of a different molecule. Usually the
electronegative atom is oxygen, nitrogen,
or fluorine.
In other words - The hydrogen on one
molecule attached to O or N that is
attracted to an O or N of a different
molecule.
5) HYDROGEN BONDING
F, O, N
The origin of hydrogen bonding
The molecules which have this extra bonding
are:
Notice that in each of these molecules:
The hydrogen is attached directly to one of the most
electronegative elements, causing the hydrogen to acquire a
significant amount of positive charge.
Each of the elements to which the hydrogen is attached is not
only significantly negative, but also has at least one "active" lone
pair.
Lone pairs at the 2-level have the electrons contained in a
relatively small volume of space which therefore has a high density
of negative charge. Lone pairs at higher levels are more diffuse and
not so attractive to positive things.
Hydrogen bonding in alcohols
It is important to realise that hydrogen bonding exists in
addition to van der Waals attractions. For example, all the
following molecules contain the same number of electrons, and
the first two are much the same length. The higher boiling point
of the butan-1-ol is due to the additional hydrogen bonding.
Comparing the two alcohols (containing -OH groups), both
boiling points are high because of the additional hydrogen
bonding due to the hydrogen attached directly to the oxygen but they aren't the same.
The boiling point of the 2-methylpropan-1-ol isn't as high as the
butan-1-ol because the branching in the molecule makes the
van der Waals attractions less effective than in the longer
butan-1-ol.
Hydrogen bonding in alcohols
An alcohol is an organic
molecule containing an -O-H
group.
Any molecule which has a
hydrogen atom attached
directly to an oxygen or a
nitrogen is capable of
hydrogen bonding. Such
molecules will always have
higher boiling points than
similarly sized molecules
which don't have an -O-H or an
-N-H group. The hydrogen
bonding makes the molecules
"stickier", and more heat is
necessary to separate them.
Ethanol, CH3CH2-O-H, and
methoxymethane, CH3-O-CH3,
both have the same molecular
formula, C2H6O.
H bonding in liquid water and ice
ice
liquid water
HYDROGEN BONDING
Hydrogen bonding is usually stronger
than normal dipole forces between
molecules. Of course hydrogen
bonding is not nearly as strong as
normal covalent bonds within a
molecule - it is only about 1/10th as
strong. It’s weaker than metallic and
ionic bonds too.
This is still strong enough to have
many important consequences on the
properties of water.
!!!!WHY?
Liquid Water has higher
density than ice?
+4°C?
Hydrogen Bonds in ice and liquid
water
In liquid water each molecule is hydrogen
bonded to approximately 3.4 other water
molecules. In ice each each molecule is
hydrogen bonded to 4 other molecules.
Compare the two structures below. Notice the
empty spaces within the ice structure
ice floats on water
GIANT COVALENT STRUCTURES
Giant covalent substances like
diamond,
graphite and
silicon dioxide (silicon(IV) oxide),
and relates those structures to the
physical properties of the substances.
The structure of diamond
The giant covalent
structure of diamond
Carbon has an
electronic
arrangement of
1s2,2s22p2. In
diamond, each carbon
shares electrons with
four other carbon
atoms - forming four
single bonds.
The structure of diamond
The diamond
crystal bond
structure gives the
gem its hardness
and differentiates it
from graphite.
The structure of diamond
In the diagram some
carbon atoms only seem
to be forming two bonds
(or even one bond), but
that's not really the case.
We are only showing a
small bit of the whole
structure.
This is a giant covalent
structure - it continues on
and on in three
dimensions. It is not a
molecule, because the
number of atoms joined
up in a real diamond is
completely variable depending on the size of
the crystal.
Diamond
The physical properties of diamond
Diamond
has a very high melting point (almost 4000°C). Very
strong carbon-carbon covalent bonds have to be broken
throughout the structure before melting occurs.
is very hard. This is again due to the need to break very
strong covalent bonds operating in 3-dimensions.
doesn't conduct electricity. All the electrons are held
tightly between the atoms, and aren't free to move.
is insoluble in water and organic solvents. There are no
possible attractions which could occur between solvent
molecules and carbon atoms which could outweigh the
attractions between the covalently bound carbon atoms.
The structure of graphite
The giant covalent
structure of graphite
Graphite has a layer
structure which is
quite difficult to draw
convincingly in three
dimensions.
The diagram on the
right shows the
arrangement of the
atoms in each layer,
and the way the layers
are spaced.
The structure of graphite
Notice that you can't really
draw the side view of the
layers to the same scale as
the atoms in the layer
without one or other part of
the diagram being either
very spread out or very
squashed.
In that case, it is important
to give some idea of the
distances involved. The
distance between the layers
is about 2.5 times the
distance between the atoms
within each layer.
The layers, of course,
extend over huge numbers
of atoms - not just the few
shown above.
You might argue that
carbon has to form 4
bonds because of its 4
unpaired electrons,
whereas in this diagram it
only seems to be forming
3 bonds to the
neighbouring carbons.
This diagram is
something of a
simplification, and shows
the arrangement of atoms
rather than the bonding.
The bonding in graphite
Each carbon atom uses three of its electrons to
form simple bonds to its three close neighbours.
That leaves a fourth electron in the bonding
level.
These "spare" electrons in each carbon atom
become delocalised over the whole of the sheet
of atoms in one layer.
They are no longer associated directly with any
particular atom or pair of atoms, but are free to
wander throughout the whole sheet.
The bonding in graphite
The bonding in graphite is like a vastly
extended version of the bonding in
benzene.
Each carbon atom undergoes sp2
hybridisation, and then the
unhybridised p orbitals on each carbon
atom overlap sideways to give a
massive pi system above and below the
plane of the sheet of atoms.
Allotropes of Carbon
Graphite to diamond animation
This animation shows how graphite becomes diamond under extreme heat and pressure
Solubility and
Intermolecular Forces
Like dissolves like
– Polar solutes dissolve in polar solvents
– Nonpolar solutes dissolve in nonpolar
solvents
Molecules with similar intermolecular
forces will mix freely
Ionic Solute with
Polar Solvent
Ionic Solute with
Nonpolar Solvent
Nonpolar Solute with
Nonpolar Solvent
Nonpolar Solute
with Polar Solvent
Comparison of the Properties of Substances with Ionic,
Covalent, Metallic or Intermolecular Bonds
Ionic
Covalent
Metallic
Intermolecular
Bond
strength
Strong
Very strong
Variable strength,
generallymoderat
e
Weak
Hardness
Moderate to high
Insulators in
solids and liquid
states
Low to moderate;
ductile, malleable
Crystal soft and
somewhat
plastic
Electrical
conductivity
Conducts by ion
transport, but only
when liquid or
dissociated
Low
Good conductors;
conducts by
electron transport
Insulators in
both solid and
liquid states
Melting
point
Moderate to high
Very high
Generally high
Low
Solubility
Soluble in polar
solvents
Very low
solubility
Insoluble except
in acids or alkalis
by chemical
reaction
Soluble in
organic
solvents
Examples
Most minerals
Diamond, oxygen,
hydrogen, organic
molecules
Cu, Ag, Au, other
metals
Organic
compounds
Boiling Point of Various Material (˚C)
Noble gas
Helium
neon
argon
He
Ne
Ar
-269
-246
-186
Nonpolar covalent
hydrogen
oxygen
methane
chlorine
H2
O2
CH4
Cl2
-253
-183
-164
-34
polar covalent
ammonia
hydrogen
fluoride
water
NH3
-33
19.5
HF
100
H2O
ionic
metallic
potassium
chloride
sodium
chloride
magnesium
oxide
KCl
771
NaCl
1413
MgO
2826
copper
iron
tungsten
Cu
Fe
W
2567
2750
5660
REFERENCES
 VIRTUAL CHEMBOOK
http://www.elmhurst.edu/%7Echm/vchembook/index.html
 Gary L. Bertrand
Department of Chemistry
University of Missouri-Rolla
 chemguideHelping you to understand Chemistry
Jim Clark 2005
http://www.chemguide.co.uk/
 Minerals
http://www.uwgb.edu/dutchs/EarthSC202Notes/minerals.htm